many-electron atoms shielded effective nuclear charge, z eff, penetrate s < p < d < f. we...
TRANSCRIPT
Many-Electron AtomsWe have to examine the balance of attractions and repulsions in the
atom to explain why subshells of a given shell have different energies. As well as being attracted by the nucleus, each electron in a many-electron atom is repelled by the other electrons present. As a result, it is less tightly bound to the nucleus than it would be if those other electrons were absent. We say that each electron is shieldedshielded from the full attraction of the nucleus by the other electrons in the
atom. The shielding effectively reduces the pull of the nucleus on an electron. The effective nuclear charge, Zeffective nuclear charge, Zeffeff,, experienced by the electron is always less than the actual nuclear charge, Z, because the electron-electron repulsions work against the pull of the nucleus. Note that the other electrons do not “block” the influence of the nucleus; they simply provide additional repulsive Coulombic interactions that partly counteract the pull of the nucleus. Finally, an s-electron of any shell can be found very close to the nucleus, so we say that it can penetratepenetrate through the inner shells. A p-electron penetrates much less. Because a p-electron penetrates less than an s-electron through the inner shells of the atom, it is more effectively shielded from the nucleus and hence experiences a smaller effective nuclear charge than an s-electron does. That is, an s-electron is bound more tightly than a p-electron and has a slightly lower (more negative) energy. In a many-electron atom, because of the effects of penetration and shielding, the order of energies of orbitals in a given shell is typically ss < < pp < < dd < < ff..
A few principles to consider:
Aufbau Principle: The procedure for arriving at the ground-state electron configurations of atoms and molecules in order of increasing atomic number. To proceed from one atom to the next, we add a proton and some neutrons to the nucleus and then describe the orbital into which the added electron goes.
Hund’s Rule: Whenever orbitals of equal energy (degenerate) are available, electrons occupy these orbitals singly before pairing begins.
Core vs. Valence electrons: inner vs. outermost electrons (latter contained within outermost shell)
Electron ConfigurationElectron configuration is a shorthand
notation for describing the arrangement of the electrons about
the nucleus.General Format using the quantum numbers:
n n l l e-e-
RULES:
1. Fill the lowest energy levels first.
1s 2s 2p 3s 3p 4s 3d 4p
2. No more than two electrons per orbital.
n = principle quantum number
l = angular momentum quantum number
e- = number of electrons
Lowest
Electron ConfigurationExamples:Examples:
HH :: 1s1 He:He: 1s2 Li :Li : 1s2 2s1
Co: 1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p6 6 4s4s22 3d 3d77
Br: 1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p6 6 4s4s22 3d 3d10 10 4p4p55
The condensed electron configuration distinguishes the core electrons from the valence electrons. CORE CORE electrons are tightly held to the nucleus and resemble a noble gas configuration. VALENCE VALENCE electrons are the outer most electrons and are involved in chemical
reactions. Examples of the condensed configuration:
Li:[He] 2s1 Co:[Ar] 4s2 3d3d7 7
Br:[Ar] 4s2 3d3d10 10 4p4p55
Electron ConfigurationThe full & condensed electron configuration for some elements:
CC 1s 1s22 2s 2s22 2p 2p22 oror [He] 2s[He] 2s22 2p 2p22
OO 1s 1s22 2s 2s22 2p 2p44 or or [He] 2s[He] 2s22 2p 2p22
NeNe 1s 1s22 2s 2s22 2p 2p66 or or [Ne][Ne]
NaNa 1s 1s22 2s 2s22 2p 2p66 3s 3s11 oror [Ne] 3s[Ne] 3s11
Si Si 1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p22 oror [Ne] 3s[Ne] 3s22 3p 3p22
ClCl 1s 1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p55 oror [Ne] 3s[Ne] 3s22 3p 3p55
ArAr 1s 1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p66 oror [Ar][Ar]
K K 1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p6 6 4s4s11 oror [Ar]4s[Ar]4s11
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s1
s2
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
s2p1 p2 p3 p4 p5
p6
f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1
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Electron Configuration fromthe Periodic Table
P = [Ne]3s23p3
P has 5 valence electrons
3p3
P
Ne1234567
1A
2A 3A 4A 5A 6A 7A
8A
3s2
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Sublevel Splitting in Multielectron Atoms
• the sublevels in each principal energy level of Hydrogen all have the same energy – we call orbitals with the same energy degenerate– or other single electron systems
• for multielectron atoms, the energies of the sublevels are split– caused by electron-electron repulsion
• the lower the value of the l quantum number, the less energy the sublevel has– s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)
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Penetrating and Shielding• the radial distribution function shows that
the 2s orbital penetrates more deeply into the 1s orbital than does the 2p
• the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they are more shielded from the attractive force of the nucleus
• the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively
• the result is that the electrons in the 2s sublevel are lower in energy than the electrons in the 2p
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Penetration & Shielding
Ene
rgy
1s
7s
2s
2p
3s
3p3d
6s6p
6d
4s
4p4d
4f
5s
5p
5d5f
Notice the following:1. because of penetration, sublevels within
an energy level are not degenerate2. penetration of the 4th and higher energy
levels is so strong that their s sublevel is lower in energy than the d sublevel of the previous energy level
3. the energy difference between levels becomes smaller for higher energy levels
QUANTUM MECHANICS & ORBITAL DIAGRAMSOrbital Energy Levels:
_ _ _ 6p _ _ _ _ _
5d _ _ _ _ _ _ ___ 4f
E 6s _ _ _n 5p _ _ _ _ _e 4dr __g 5s _ _ _y 4p _ _ _ _ _
3d__4s _ _ _
3p__3s _ _ _
2p__2s__
1s
Example of Ionization Energies:
Al(g) Al+(g) + e- I1 = 580 kJ/mol
Al+(g) Al2+
(g) + e- I2 = 1815 kJ/molAl2+
(g) Al3+(g) + e- I3 = 2740 kJ/mol
Al3+(g) Al4+
(g) + e- I4 = 11,600 kJ/mol
Orbital DiagramsOrbital DiagramsOrbital diagrams are written in order of increasing energy
levels starting with the lowest energy level the 1s orbital.
___ ___ ___ 4p
___ ___ ___ ___ ___ 3d
___ 4s
___ ___ ___ 3p
___ 3s
___ ___ ___ 2p
___ 2s
___ 1s
RULES:
(1) fill the lowest energy level first
(2) fill each orbital in a subshell with one electron first before you double up.
(3) Completely fill each subshell before proceeding to the next energy level.
Remember the Remember the order!!order!!
Orbital DiagramsOrbital DiagramsFill in the orbital diagrams for:
C C OO
___ ___ ___ 4p ___ ___ ___ 4p
__ __ __ __ __ 3d __ __ __ __ __ 3d
___ 4s ___ 4s
___ ___ ___ 3p ___ ___ ___ 3p
___ 3s ___ 3s
___ ___ ___ 2p ___ ___ ___ 2p
___ 2s ___ 2s
___ 1s ___ 1s
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Transition Elements• for the d block metals, the principal energy level is one less than
valence shell– one less than the Period number– sometimes s electron “promoted” to d sublevel
4s 3d
ZnZ = 30, Period 4, Group 2B[Ar]4s23d10
• for the f block metals, the principal energy level is two less than valence shell two less than the Period number they really belong tosometimes d electron in configurationEuZ = 63, Period 6[Xe]6s24f 7 6s 4f
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Electron Configuration fromthe Periodic Table
As = [Ar]4s23d104p3
As has 5 valence electrons
As
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1A
2A 3A 4A 5A 6A 7A
8A
4s2
Ar3d10
4p3
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Practice – Use the Periodic Table to write the short electron configuration and orbital diagram for each
of the following
• Na (at. no. 11)
• Te (at. no. 52)
• Tc (at. no. 43)
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Practice – Use the Periodic Table to write the short electron configuration and orbital diagram for each
of the following
• Na (at. no. 11) [Ne]3s1
• Te (at. no. 52) [Kr]5s24d105p4
• Tc (at. no. 43) [Kr]5s24d5
3s
5s 5p4d
5s 4d
1. Determine the ground-state electron configuration for each of the following elements:
A. sulfur B. polonium
2. Predict the number of valence electrons present in each of the following atoms (include the outermost d-electrons when necessary):
A. B B. Ba C. Bi
3. Determine the ground-state electron configuration for each of the following ions:
A. Al+3 B. Tc+4
4. Predict the number of valence electrons present for each of the following ions:
A. In+ B. Tc+2
5. Give the ground-state electron configuration and number of unpaired electrons expected for each of the following ions:
A. Ga3+ B. Cu+2
6. For each of the following ground-state ions, predict the type of orbital that the electrons of highest energy will occupy:
A. Fe+2 B. Bi+3
Lecture Questions
Workshop on electron configuration1. Determine the ground-state electron configuration for each of the following elements (see last page of this section for sample energy levels):
A. chlorine B. cesiumC. vanadium D. rhenium
2. Predict the number of valence electrons present in each of the following atoms (include the outermost d-electrons):
A. Sn B. La C. Mn D. Zn3. Determine the ground-state electron configuration for each of the following ions:
A. Co+3B. Mo+2 C. Ra+2
D. I- E. Ir+ F. Ru+4
4. Predict the number of valence electrons present for each of the following ions:A. Tl+ B. Po+2 C. Ta+2 D. Re+
5. Give the ground-state electron configuration and number of unpaired electrons expected for each of the following ions:
A. Ga+ B. Cu+1C. Pb+2 D. Se-2
6. For each of the following ground-state ions, predict the type of orbital that the electrons of highest energy will occupy:
A. Fe+3 B. B+3 C. As+3 D. Os+
PERIODICITYDiamagnetic vs. Paramagnetic species:
Diamagnetic has all its electrons paired and is slightly repelled by a magnetic field
Paramagnetic has one or more unpaired electrons and is attracted into a magnetic field.
Which group(s) on the periodic table will have elements that are always diamagnetic?
Periodic Trends1. Atomic Radius
2. Ionization Energy – energy needed to remove an electron from gaseous atom
Decreases acrossincreases
Increases down
Increases acrossDecreases down
3. Electron Affinity – energy released when an electron is added to gaseous atom
4. Electronegativity – the electron pulling power of an atom when it is part of a molecule (denoted with the Greek letter )
5. Metallic Character
Workshop on periodic trends1. Arrange the following in terms of DECREASING atomic radius & then first ionization energy & then electronegativity:
Be, B, C, N, O, F, Ne
2. Why is the first ionization energy of aluminum slightly lower than the first ionization energy for magnesium?
3. Why is the second ionization energy for sodium so much greater than its first ionization energy?
4. Arrange the following in terms of DECREASING atomic (or ionic) radii: O+, O, O-
5. Give a reason why the electronegativity for F is so much greater than the electronegativity for Fr.